Deck 16: Reactions Between Acids and Bases

Question 1

In the titration of 0.100 M HCl with the titrant 0.100 M NaOH, what species are present at any point after NaOH has been added but before the equivalence point?

A) HCl only
B) NaOH only
C) NaCl only
D) HCl and NaOH
E) HCl and NaCl

Answer 1

HCl and NaCl

Question 2

In the titration of 0.100 M HCl with the titrant 0.100 M NaOH, what species are present at the equivalence point?

A) HCl only
B) NaOH only
C) NaCl only
D) NaOH and NaCl
E) HCl and NaCl

Answer 2

NaCl only

Question 3

Calculate the pH of a solution made by mixing 20.00 mL of 0.300 M HCl and 50.00 mL of 0.100 M NaOH and diluting to a final volume of 1.00 L.

A) 3.00
B) 2.22
C) 2.30
D) - 1.18
E) 0.52

Answer 3

3.00

Question 4

Calculate the volume of 0.100 M HCl required to neutralize 1.00 g of Ba(OH)₂ (molar mass = 171.3 g/mol).

A) 58.4 mL
B) 117 mL
C) 29.2 mL
D) 585 mL
E) no reaction occurs

Question 5

Calculate the pH of a solution after 10.0 mL of 0.100 M NaOH is added tO₄0.0 mL of 0.250 M HBr.

A) 0.65
B) 1.00
C) - 1.0
D) 13.01
E) 0.74

Question 6

A 25.0 mL portion of 0.0500 M HCl was added to 10.0 mL of 0.0250 M NaOH. Calculate the pH of the resulting solution.

A) 12.54
B) 1.54
C) 3.00
D) 12.06
E) 1.30

Question 7

What volume of 0.200 M H₂ SO₄is required to neutralize 50.0 mL of 0.100 M NaOH?
Assume the chemical equation is:
H₂ SO₄(aq) + 2 NaOH (aq) Na₂ SO₄(aq) + 2 H₂ O ( )

A) 50.0 mL
B) 25.0 mL
C) 100 mL
D) 12.5 mL
E) 6.25 mL

Question 8

Solutions are made by combining equal amounts of the following. Which is a buffer?

A) 0.1 M HCl + 0.1 M Cl ^-
B) 0.1 M HF + 0.1 M Na⁺
C) 0.1 M HF + 0.1 M NaOH
D) 0.1 M NaF + 0.05 M HF
E) 0.1 M NaF + 0.05 M Na⁺

Question 9

What is the molarity of Ba(OH)₂ solution if 65.0 mL of 0.0500 M HCl is required to titrate 35.00 mL of the Ba(OH)₂ solution?

A) 3.25×10 ^{- 3} M
B) 9.28×10 ^{- 2} M
C) 4.64×10 ^{- 2} M
D) 2.69×10 ^{- 3} M
E) 1.69×10 ^{- 2} M

Question 10

In the titration of 50.0 mL of a 0.100 M HCl solution with 0.100 M NaOH, what is the pH of the solution after the addition of 20.0 mL of the NaOH solution?

A) 1.00
B) 1.22
C) 1.15
D) 1.37
E) 12.63

Question 11

Which pair(s) of substances would make a suitable buffer pair combination?
I. NaOH and NaBr
II. HF and NaF
III. CH₃NH₃⁺Cl ^- and CH₃NH₂

A) I only
B) II only
C) III only
D) I and II
E) II and III

Question 12

Calculate the pH of a solution formed when 1.5 L oF₄.5×10 ^{- 3} M HCl is mixed witH₂ .0 L of 3.0×10 ^{- 3} M HNO₃.

A) 3.67
B) 3.74
C) 2.25
D) 3.16
E) 2.44

Question 13

Which of the following general combinations of substances form an acid-base buffer when added to water?

A) a strong acid and the sodium salt of its conjugate base
B) a strong base and the chloride salt of its conjugate acid
C) a weak acid and the sodium salt of its conjugate base
D) a weak acid
E) a weak base and a strong base

Question 14

During a titration, 60.0 mL of 0.010 M NaOH is added to 35.0 mL of 0.020 M HCl solution. What is the pH of the resulting solution?

A) 1.70
B) 12.00
C) 2.98
D) 5.41
E) 4.05

Question 15

If it takes 32.0 mL of 0.100 M HCl to titrate 25.0 mL of a Ba(OH)₂ solution to the equivalence point, what is the molarity of the original Ba(OH)₂ solution?

A) 0.128 M
B) 3.20×10 ^{- 2} M
C) 0.256 M
D) 6.40×10 ^{- 2} M
E) 0.100 M

Question 16

A 350 mL volume of 0.150 M Ba(OH)₂ was completely neutralized by 35.0 mL of HCl. What is the molarity of the HCl solution?

A) 1.50 M
B) 3.00 M
C) 0.750 M
D) 1.00 M
E) 0.500 M

Question 17

What is the pH of a solution of 25.0 mL of 0.250 M HNO₃ after the addition oF₄0.0 mL of 0.300 M KOH?

A) 1.02
B) 0.60
C) 12.95
D) 13.25
E) 13.45

Question 18

In the titration of 0.100 M HCl with the titrant 0.100 M NaOH, what species are present after the equivalence point?

A) HCl only
B) NaOH only
C) NaCl only
D) NaOH and NaCl
E) HCl and NaCl

Question 19

A 0.40 gram sample of NaOH (molar mass = 40.0 g/mol) was completely neutralized witH₆2.5 mL of H₂ SO₄solution. Assume the chemical equation is:
H₂ SO₄(aq) + 2 NaOH (aq) Na₂ SO₄(aq) + 2 H₂ O ( ) The molarity of the H₂ SO₄solution was:

A) 0.080 M
B) 0.16 M
C) 0.62 M
D) 6.4 M
E) 3.2 M

Question 20

What is the volume of 0.0175 M Ba(OH)₂ required to neutralize 10.0 mL of 0.0300 M HCl?

A) 10.0 mL
B) 34.2 mL
C) 17.1 mL
D) 8.6 mL
E) 4.3 mL

Question 21

Which one of these combinations would give a buffer that would be most effective in the pH range 3.0 tO₄.0?

A) 0.10 M CH₃COOH, 0.10 M CH₃COONa, K _a = 1.8×10 ^{- 5}
B) 0.10 M NH₄Cl, 0.10 M NH₃, K _b = 1.8×10 ^{- 5}
C) 0.10 M HF, 0.10 M NaF, K _a = 3.54×10 ^{- 4}
D) 0.10 M H₂CO₃, 0.10 M NaHCO₃, K _a = 4.30×10 ^{- 7}
E) 0.10 M HCl, 0.10 NaCl

Question 22

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions is the pH greater than the p K _a of the weak acid?

A) When the p K _a equals 7.0.
B) When [HX] > [NaX].
C) When [HX] < [NaX].
D) When [HX] = [NaX].
E) When there is no NaX present.

Question 23

What is the pH of a buffer solution that consists of 0.83 M HBrO and 0.53 M KBrO?
K _a(HBrO) = 2.3×10 ^{- 9}

A) pH = 8.19
B) pH = 8.45
C) pH = 8.83
D) pH = 8.93
E) pH = 12.05

Question 24

A solution of HCOOH ( K _a = 1.8×10 ^{- 4}) and HCOO ^- was submitted for chemical analysis. The results were:
[HCOOH] = 0.050 M, [HCOO ^- ] = 0.15 M. Calculate the pH.

A) 4.22
B) 3.74
C) 3.27
D) 1.30
E) - 1.30

Question 25

Consider the buffer pair, NH₄Cl/NH₃. What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted to 12.00 and the p K _a for NH₄⁺ equals 9.25?

A) This buffer solution contains relatively more NH₄⁺ than NH₃.
B) This buffer solution contains relatively more NH₃ than NH₄⁺.
C) This buffer solution contains equal concentrations of NH₄⁺ and NH₃.
D) More information is needed to determine the relative proportion of each member.

Question 26

Which pair(s) of substances would make a suitable buffer pair ?
I. HCl and NaCl
II. HF and NaF
III. NH₄Cl and NH₃

A) I only
B) II only
C) III only
D) I and II
E) II and III

Question 27

A buffer solution maintains a constant pH level to a certain extent when a strong base such as NaOH is added. Which of the following buffer solutions has the greatest buffering capacity to consume added NaOH?

A) 0.5 M HF/1.0 M NaF
B) 0.5 M HF/1.5 M NaF
C) 1.0 M HF/0.5 M NaF
D) 1.0 M HF/1.5 M NaF
E) 1.5 M HF/0.5 M NaF

Question 28

What is the pH of a solution that is 0.100 M methylamine (CH₃NH₂) and 0.200 M methyl ammonium chloride (CH₃NH₃Cl)?
The K _b of methylamine is 3.70×10 ^{- 4}.

A) 1.00
B) 3.73
C) 10.73
D) 3.43
E) 10.27

Question 29

Which of the following would not make a suitable buffer pair?

A) NH₄Cl/NH₃
B) HNO₃/NaNO₃
C) NaHCO₃/Na₂CO₃
D) H₃PO₄/KH₂PO₄
E) HCN/NaCN

Question 30

Consider the two buffer solutions prepared below:
Solution A:
[HA] = 0.15 M and [A ^- ] = 0.25 M Solution B:
[HA] = 1.50 M and [A ^- ] = 2.50 M Which statement below is true regarding the pH and buffering capacity of these two solutions?

A) Solution A will have a greater pH value and a smaller buffering capacity.
B) Solution B will have a greater pH value and a smaller buffering capacity.
C) Solution B will have a greater pH value and a larger buffering capacity.
D) The pH of both solutions will be identical but Solution A will have a larger buffering capacity.
E) The pH of both solutions will be identical but Solution B will have a larger buffering capacity.

Question 31

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions does the pH just equal the p K _a of the weak acid?

A) When the p K _a equals 7.0.
B) When [HX] > [NaX].
C) When [HX] < [NaX].
D) When [HX] = [NaX].
E) When there is no NaX present.

Question 32

Calculate the pH of a solution that is 0.78 M NH₃ and 0.140 M NH₄NO₃. K _b = 1.8×10 ^{- 5} for NH₃.

A) 4.00
B) 10.00
C) 4.74
D) 9.26
E) 13.91

Question 33

Calculate the pH of a buffer prepared by dissolving 0.10 mol of NH₄Cl in 1.00 L of 0.15 M NH₃. K _b = 1.8×10 ^{- 5} for NH₃.

A) 9.43
B) 4.56
C) 11.34
D) 2.72
E) 4.74

Question 34

What is the pH of an aqueous solution that is 0.55 M HNO₂ and 0.75 M KNO₂?
K _a(HNO₂) = 7.1×10 ^{- 4}

A) pH = 3.01
B) pH = 3.28
C) pH = 3.63
D) pH = 6.94
E) pH = 7.56

Question 35

What is the pH of an aqueous solution that contains 0.085 M HNO₂ and 0.10 M potassium nitrite (KNO₂)?
K _a(HNO₂) = 4.5×10 ^{- 4}

A) pH = 3.27
B) pH = 3.42
C) pH = 4.28
D) pH = 4.42
E) pH = 7.87

Question 36

Consider the buffer pair, HF/F ^- . What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted tO₄.00 and the p K _a for HF equals 3.17?

A) This buffer solution contains relatively more HF than F ^- .
B) This buffer solution contains relatively more F ^- than HF.
C) This buffer solution contains equal concentrations of HF and F ^- .
D) More information is needed to determine the relative proportion of each member.

Question 37

A solution is made by dissolving 0.100 mole NH₃ ( K _b = 1.8×10 ^{- 5}) and 0.200 mole of NH₄Cl in water and diluting to 1.00 L. What is the pH of the resulting solution?

A) 8.95
B) 11.12
C) 2.72
D) 13.00
E) 5.04

Question 38

What is the pH of a solution prepared by mixing 0.250 mol of hydrazoic acid, HN₃, and 0.500 mol of sodium azide, NaN₃, to make 1.00 liter of solution?
( K _a = 1.9×10 ^{- 5})

A) 4.42
B) 5.02
C) 9.58
D) 4.72
E) 8.98

Question 39

Which of the following mixtures would not be suitable for preparing a buffer solution?

A) NH₄Cl and NH₃
B) HNO₃ and NaNO₃
C) HF and KF
D) C₆H₅COOH (Benzoic acid) and C₆H₅COO ^- Na⁺
E) CH₃COOH and CH₃COO ^- Na⁺

Question 40

A buffer solution maintains a constant pH level to a certain extent when a strong acid such as HCl is added. Which of the following buffer solutions has the greatest buffering capacity to consume added HCl?

A) 0.5 M HC₂H₃O₂/0.5 M NaC₂H₃O₂
B) 1.0 M HC₂H₃O₂/0.5 M NaC₂H₃O₂
C) 1.0 M HC₂H₃O₂/1.5 M NaC₂H₃O₂
D) 1.5 M HC₂H₃O₂/1.0 M NaC₂H₃O₂
E) 1.5 M HC₂H₃O₂/0.5 M NaC₂H₃O₂

Question 41

What is the pH after the addition of 0.20 mole of HCl to 1.00 liter of 1.00 M CH₃COONa and 1.00 M CH₃COOH ( K _a = 1.8×10 ^{- 5}) buffer solution?
(Assume no volume change.)

A) 0.70
B) 2.31
C) 4.57
D) 9.28
E) 13.3

Question 42

A buffer was prepared by mixing 0.60 moles of HF and 0.40 moles of NaF in enough water to make 1.00 L of solution. What is the pH of this solution?
K _a = 3.5×10 ^{- 4} for HF.

A) 10.72
B) 3.48
C) 3.58
D) 10.42
E) 3.28

Question 43

Calculate the amount of ammonium chloride (NH₄Cl) that must be added to 1.0 L of 0.10 M NH₃ ( K _b = 1.8×10 ^{- 5}) to prepare a buffer of pH 9.00.

A) 0.10 mole
B) 0.056 mole
C) 0.18 mole
D) 5.6×10 ^{- 5} mole
E) buffer cannot be made

Question 44

A 0.10 M solution of hydrazine (N₂H₄, K _b = 1.7×10 ^{- 6}) containing an unknown concentration of hydrazine hydrochloride (N₂H₅Cl) has a pH of 7.50. What is the concentration of N₂H₅Cl in the solution?

A) 1.26 M
B) 0.54 M
C) 5.8 M
D) 0.10 M
E) 0.72 M

Question 45

How many grams of sodium acetate, CH₃COONa, is required to mix witH₂ .00 Liters of 0.450 M acetic acid, CH₃COOH, in order to prepare a buffer solution with a pH that equals 5.00?
P K _a(CH₃COOH) = 4.74 and MM(CH₃COONa) = 82.03 g/mol

A) 2.00×10 ^{- 2} g CH₃COONa
B) 45.6 g CH₃COONa
C) 67.2 g CH₃COONa
D) 134 g CH₃COONa
E) 299 g CH₃COONa

Question 46

A buffer was prepared such that the pH was 5.20. If the solution contains 0.400 moles of acetic acid (CH₃COOH) per liter, what is the concentration of the acetate ion (CH₃COO ^- ) in this solution?
The p K _a for acetic acid is 4.74.

A) 0.40 M
B) 0.14 M
C) 4.2×10 ^{- 3} M
D) 6.3×10 ^{- 6} M
E) 1.1 M

Question 47

Exhibit 16-1 Consider the equilibrium reaction below to answer the following question(s). HNO₂ + H₂ O H₃O + + NO₂ 1 - K a (HNO₂ ) = 4.5 10 - 4
Refer to Exhibit 16-1. What is the concentration of nitrite ion, NO₂ - , present at equilibrium if 0.100 moles of HCl are added to 0.200 moles of HNO₂ to form one-liter of solution?

A) [NO₂ ^- ] = 9.00×10 ^{- 4} M
B) [NO₂ ^- ] = 2.12×10 ^{- 3} M
C) [NO₂ ^- ] = 9.49×10 ^{- 3} M
D) [NO₂ ^- ] = 0.100 M
E) [NO₂ ^- ] = 0.200 M

Question 48

What is the pH of a solution that consists of 0.25 M HC₂H₃O₂ and 0.35 M NaC₂H₃O₂?
K _a(HC₂H₃O₂) = 1.8×10 ^{- 5}

A) 0.60
B) 4.60
C) 4.89
D) 5.50
E) 6.07

Question 49

What is the concentration of ammonium ion, [NH₄⁺], present at equilibrium if 0.100 moles of NaOH are added to 0.200 moles of NH₃ to form one-liter of solution?

A) [NH₄⁺] = 9.0×10 ^{- 6} M
B) [NH₄⁺] = 3.6×10 ^{- 5} M
C) [NH₄⁺] = 1.9×10 ^{- 3} M
D) [NH₄⁺] = 6.0×10 ^{- 3} M
E) [NH₄⁺] = 0.10 M

Question 50

Which of the following titration curves listed below best represents a curve for the complete titration of oxalic acid, H₂C₂O₄ (a diprotic acid), with a strong base such as NaOH?

A)
B)
C)
D)
E)

Question 51

Consider the ionization of acetic acid as follows:CH ₃COOH + H₂ O H₃O + + CH₃COO 1 - What is the effect of adding the common ion, acetate, CH₃COO - (source:
CH₃COONa), to this equilibrium system?

A) The equilibrium shifts to the left and the pH rises.
B) The equilibrium shifts to the right and the pH rises.
C) The equilibrium shifts to the left and the pH lowers.
D) The equilibrium shifts to the right and the pH lowers.
E) Nothing happens. The pH remains constant and the equilibrium does not shift in either direction.

Question 52

What is the pH of a buffer solution that consists of 0.25 M HClO₂ and 0.75 M KClO₂?
K _a(HClO₂) = 1.1×10 ^{- 2}

A) pH = 1.48
B) pH = 1.96
C) pH = 2.44
D) pH = 3.41
E) pH = 5.61

Question 53

What is the pH after the addition of 0.050 moles of NaOH to 1.00 L of a 0.500 M NH₃ / 0.500 M NH₄Cl buffer solution?
Assume that there is no volume change on addition of the NaOH. K _b for NH₃ = 1.8×10 ^{- 5}.

A) 4.74
B) 4.66
C) 9.34
D) 9.17
E) 9.53

Question 54

A buffer is made by dissolving 0.10 mol of NaF in 1.00 L of 0.20 M HF. What is the pH of this buffer?
K _a(HF) = 6.3×10 ^{- 4}

A) 1.05
B) 1.58
C) 3.16
D) 3.46
E) 3.70

Question 55

Consider the ionization of hydrofluoric acid as follows:
HF + H₂ O H₃O + + F 1 - What is the effect of adding the common ion, F - (source:
NaF), to this equilibrium system?

A) The equilibrium shifts to the right and the pH lowers.
B) The equilibrium shifts to the left and the pH lowers.
C) The equilibrium shifts to the right and the pH rises.
D) The equilibrium shifts to the left and the pH rises.
E) Nothing happens. The pH remains constant and the equilibrium does not shift in either direction.

Question 56

What is the pH of an aqueous solution that contains 0.183 M sodium formate (NaCHO₂) and 0.300 M in formic acid?
K _a(HCHO₂) = 1.8×10 ^{- 4}

A) pH = 3.53
B) pH = 3.96
C) pH = 4.56
D) pH = 5.16
E) pH = 8.13

Question 57

Exhibit 16-1 Consider the equilibrium reaction below to answer the following question(s). HNO₂ + H₂ O H₃O + + NO₂ 1 - K a (HNO₂ ) = 4.5 10 - 4
Refer to Exhibit 16-1. What is the effect of adding NaNO₂ to an aqueous solution of HNO₂ as shown in the equilibrium reaction above?

A) The equilibrium will shift left favoring the reactant side and the pH will drop.
B) The equilibrium will shift right favoring the product side and the pH will drop.
C) There will be no shift in the equilibrium reaction and the pH will remain constant.
D) The equilibrium will shift left favoring the reactant side and the pH will rise.
E) The equilibrium will shift right favoring the product side and the pH will rise.

Question 58

What is the pH of an aqueous solution that contains 0.39 M C₆H₅COOH and 0.14 M C₆H₅COONa?
K _a(C₆H₅COOH) = 6.3×10 ^{- 5}

A) pH = 0.409
B) pH = 3.76
C) pH = 4.65
D) pH = 6.29
E) pH = 8.65

Question 59

Consider the following equilibrium reaction:
NH₃+ H₂ O NH ₄+ + OH - K b (NH ₃) = 1.8 10 - 5 Qualitatively, what is the effect of adding NaOH to an aqueous solution of NH₃as shown in the equilibrium reaction above?
(Use Le Chatelier s principle.)

A) The equilibrium will shift left favoring the reactant side and the pH will drop.
B) The equilibrium will shift right favoring the product side and the pH will drop.
C) There will be no shift in the equilibrium reaction and the pH will remain constant.
D) The equilibrium will shift left favoring the reactant side and the pH will rise.
E) The equilibrium will shift right favoring the product side and the pH will rise.

Question 60

Calculate the number of moles of sodium acetate (CH₃COONa) that must be added to 1.00 L of 0.100 M CH₃COOH ( K _a = 1.8×10 ^{- 5}) to give a buffer with pH₅.00.

A) 1.0×10 ^{- 5}
B) 0.18
C) 0.057
D) 0.10
E) 0.20

Question 61

The titration of a weak acid, HA ( K _a = 1.0×10 ^{- 4}), with NaOH requires 40.0 mL to reach the equivalence point. The pH after 20 mL is

A) Strongly alkaline, approximately 12
B) Moderately alkaline, approximately 10
C) Slightly alkaline, approximately 8
D) Neutral, approximately 7
E) Moderately acidic, approximately 4

Question 62

Which of the following titration curves listed below best represents a curve for the complete titration of the weakly basic dianion, sulfide, S^{2 -} , with a strong acid such as HCl as shown below in the net ionic equation?
S^{2 -} (aq) + 2 HCl (aq)→H₂S (g) + 2 Cl ^-

A)
B)
C)
D)
E)

Question 63

The pH at the equivalence point in the titration of 0.10 M NH₃ with 0.10 M HCl is:

A) strongly acidic.
B) slightly acidic.
C) neutral (pH 7.0).
D) slightly basic.
E) strongly basic.

Question 64

Exhibit 16-3 Consider the Titration curve below for the titration of a weak acid ,CH ₃COOH , with a strong base , NaOH, to answer the following question(s). CH ₃COOH + NaOH CH ₃COONa + H₂ O
Refer to Exhibit 16-3. What approach listed below would be used to solve for the pH at the equivalence point on this titration curve?

A) The pH at the equivalence point for a weak acid equals 7.0.
B) The pH equals the p K _a of the weak acid at the equivalence point.
C) The pH at the equivalence point can be determined by substituting the original molar concentration of the weak acid into the equation: pH = - log[HA].
D) The Henderson-Hasselbalch equation can be used to solve for the pH at the equivalence point.
E) The pH at the equivalence point must be solved as a problem of equilibrium involving a basic salt.

Question 65

Which of the following titrations would produce a pH equal to 7.0 at the equivalence point?
I. Titrating NH₃ with standard HCl (delivered from the burette)
II. Titrating HCl with standard NaOH (delivered from the burette)
III. Titrating CH₃COOH with standard NaOH (delivered from the burette)

A) I only
B) II only
C) III only
D) II and III
E) All of these

Question 66

Consider titrating a triprotic acid with standard NaOH as shown in the titration curve below. What is the relationship between the first equivalence point and the third equivalence point in this titration curve?

A) The pH of the first equivalence point is three times larger than the pH found at the third equivalence point.
B) The pH of the first equivalence point is three times smaller than the pH found at the third equivalence point.
C) The volume of titrant required to reach the third equivalence point is three times as large as the volume of titrant required to reach the first equivalence point.
D) The volume of titrant required to reach the first equivalence point is three times as large as the volume of titrant required to reach the third equivalence point.
E) There are no relationships between the first and third equivalence points.

Question 67

The titration of a weak acid, HA ( K _a = 1.0×10 ^{- 4}), with NaOH requires 40.0 mL to reach the equivalence point. In order to calculate the pH of the system after 20 mL of NaOH is added, the best strategy is to recognize that the system is

A) a strong acid
B) a weak acid
C) a buffer
D) neutral
E) a weak base

Question 68

In the titration of 25.0 mL of 0.100 M acetic acid ( K _a = 1.8×10 ^{- 5}) with 0.100 M NaOH, the pH after adding 12.5 mL of titrant is:

A) 12.52
B) 3.50
C) 7.00
D) 4.74
E) 2.38

Question 69

Consider the titration of a weak acid with a strong base . Which statement below is true regarding the equivalence point in this titration?

A) It occurs at a pH = 7.0 because the acid has been neutralized.
B) It occurs at a pH less than 7.0 because an acidic salt is formed.
C) It occurs at a pH less than 7.0 because a basic salt is formed.
D) It occurs at a pH greater than 7.0 because an acidic salt is formed.
E) It occurs at a pH greater than 7.0 because a basic salt is formed.

Question 70

A 20.0 mL sample of lactic acid (monoprotic with K _a = 1.37×10 ^{- 4}) requires 30.0 mL of 0.200 M NaOH for titration to the equivalence point. What is the concentration of the lactic acid solution?

A) 0.120 M
B) 0.200 M
C) 0.300 M
D) 6.00×10 ^{- 3} M
E) cannot be found without the molar mass of lactic acid.

Question 71

Which of the following titration curves listed below best represents a curve for the complete titration of citric acid, H₃C₆H₅O₇, a triprotic acid with a strong base such as NaOH?

A)
B)
C)
D)
E)

Question 72

The titration of a weak base ( K _b = 1.0×10 ^{- 4}) with HCl requires 40.0 mL to reach the equivalence point. In order to calculate the pH of the system after 20 mL of HCl is added, the best strategy is to recognize that the system is

A) a strong base
B) a weak base
C) a buffer
D) neutral
E) a strong acid

Question 73

A sample of 0.100 mole of acetic acid ( K _a = 1.8×10 ^{- 5}) in 50.0 mL of solution is titrated with standard 0.100 M NaOH. What is the pH in the titration flask after addition of 25.0 mL of the base?

A) 2.88
B) 3.18
C) 2.60
D) 4.74
E) 11.40

Question 74

Exhibit 16-3 Consider the Titration curve below for the titration of a weak acid ,CH ₃COOH , with a strong base , NaOH, to answer the following question(s). CH ₃COOH + NaOH CH ₃COONa + H₂ O
Refer to Exhibit 16-3. What approach listed below would be used to determine the pH at point A of the solution?

A) This is a problem of solving for the pH of an aqueous weak acid solution.
B) This is a problem of solving for the pH of an aqueous weak base solution.
C) This is a problem of solving for the pH of an aqueous buffer solution.
D) This is a problem of solving for the pH of an aqueous strong base solution.
E) This is a problem of solving for the pH of an aqueous salt solution.

Question 75

Exhibit 16-2 Consider titrating CH₃COOH with standard NaOH (delivered from the burette) to answer the following question(s).
Refer to Exhibit 16-2. How would the pH be determined for the aqueous solution in the titration above at the equivalence point ?

A) The pH would equal 7.0 at the equivalence point.
B) The pH is determined by the Henderson-Hasselbalch equation.
C) The pH is determined by taking - log[CH₃COOH]_original.
D) The pH is determined by solving for the pH of a basic salt.
E) The pH is determined by solving for the pH of an acidic salt.

Question 76

What is the pH of solution formed by adding 25.0 mL of 0.100 M HCl to 25.0 mL of 0.100 M pyridine (C₅H₅N, K _b = 1.8×10 ^{- 9})?

A) 2.38
B) 2.52
C) 2.98
D) 3.28
E) 8.74

Question 77

What is the pH of the following solution?
20)0 mL of 0.50 M acetic acid ( K _a = 1.8×10 ^{- 5}) is added tO₅ .00 mL of 0.50 M NaOH.

A) 4.27
B) 4.74
C) 5.22
D) 5.40
E) 7.00

Question 78

The sharpest inflection point is observed in the titration of

A) 0.1 M HCl with 0.10 M NaOH
B) 0.01 M HCl with 0.10 M NaOH
C) 0.1 M KOH with 0.10 M NaOH
D) 0.1 M HF ( K_a = 3.5×10 ^{- 4}) with 0.10 M NaOH
E) 0.10 M NH₃ with 0.10 M NaOH

Question 79

Which experiment listed below would provide a titration curve that resembles the titration shown below?

A) Titrating NH₃ with standard NaOH (delivered from the burette).
B) Titrating NH₃ with standard HCl (delivered from the burette).
C) Titrating H₃PO₄ with standard NaOH (delivered from the burette).
D) Titrating Na₃PO₄ with standard HCl (delivered from the burette).
E) Titrating Na₃PO₄ with standard NaOH (delivered from the burette).

Question 80

Exhibit 16-2 Consider titrating CH₃COOH with standard NaOH (delivered from the burette) to answer the following question(s).
Refer to Exhibit 16-2. What compound(s) is(are) present in the region after the titration has begun but before the equivalence point?
I. CH₃COOH
II. NaOH
III. CH₃COO ^- Na⁺

A) I only
B) I and II
C) I and III
D) II and III
E) All of these are present.